Higher tier

Le Chatelier's principle

Portrait of the French chemist Henri Le Chatelier.

Henry Le Chatelier was a pioneering French chemist born in 1850, whose work and research have enable scientists to better understand chemical equilibrium. His most famous contribution is perhaps the principle that bears his name, Le Chatelier's Principle. This is a theory to help chemists and other scientists understanding how chemical systems (the reacting chemicals) at equilibrium respond to changes in the reaction conditions, whether it be temperature, pressure, or concentration. In essence, this principle states that a chemical system will change its composition to minimize the effect of any changes in temperature pressure or concentration that are made to any of the substances involved in the equilibrium reaction.

Now recall that equilibrium is the point in a reversible reaction where the rate of the forward (R_f) reaction and reverse reaction (R_b) are the same. We can show this as: During a reversible reaction the reactants turn into products and the products back into reactants again. The important point to remember is that these two reactions are occurring at the same time. The forward reaction will proceed at a given rate, labelled R_f and the back reaction which turns products back into reactants will also proceed at a given rate, labelled R_b in the example below.

When a chemical reaction starts usually the rate of the forward reaction is very fast; this is simply because there are lots of reactant particles around to collide and react with each other. However as the reaction proceeds the rate of the forward reaction (R_f) will start to slow down as the amount of reactants particles available to react reduces. The opposite is true for the back reaction; to begin with there will not be many products particles available to react with each other but as the reaction proceeds the amounts of products present will gradually increase and the rate of the back reaction (R_b) will increase. Eventually after a certain period of time the rate of the forward and reverse reactions will be the same. At this point the reversible reaction has reached dynamic equilibrium.

Imagine looking a reversible chemical reaction where the rates of the forward and reverse reactions were the same; what do you think you would see happening? Well as quickly as the reactants are turning into products the products are turning back into reactants, so the amounts of the reactants and products will not be changing. The reaction might appear as if it has stopped but it has NOT. The reactants are turning into products and the products are turning back into reactants at the same rate; the reaction most definitely has not stopped; it has reached a balance point where the rates of the forward and reverse reactions are the same. As a chemist we would say the reaction has reached or achieved dynamic equilibrium. The word equilibrium implies balance and dynamic implies movement. The reaction has reached equilibrium because R_f=R_b but the reaction is still occurring, its dynamic!

Altering the position of equilibrium

From the work you have done on rates of reaction you probably already know that it is possible to change the rate of a reaction by say heating it up or increasing the concentration of a particular reactant, adding a catalyst or increasing the surface area of a reacting substance. A common misconception that a lot of students have is that if a reaction has achieved dynamic equilibrium then it consists of a 50:50 mixture of reactants and products; however this is rarely the case.

Equilibrium is a very stable low energy point for a reaction; at equilibrium the reaction is nice and happy!! It's like the ball at the bottom of the curve. If you push on it so that it moves away from the bottom of the curve the ball will roll back down again; to the low energy point. Well chemical reactions are a bit like that!- if a reaction is at equilibrium (rate of forward and reverse reactions are the same) and you come along and heat it up or put it under pressure and generally "annoy it" then the reaction will re-adjust itself to get back to a new equilibrium position, that is a new low energy state. At this new equilibrium position the amounts of reactants and products maybe different but the rate of the forward and reverse reactions will be the same e.g. consider the reaction of nitrogen and hydrogen to make ammonia; the Haber process.

Nitrogen_(g) + hydrogen_(g) ⇌ ammonia_(g)

N_2(g) + 3H_2(g) ⇌ 2NH_3(g)

At equilibrium the mixture contains only a small amount of ammonia and is mostly nitrogen and hydrogen. We would say the position of equilibrium lies to the left, that is on the reactants side. Since ammonia is a useful substance we need to shift the position of equilibrium so that it moves more to the right, that is the equilibrium mixture has a larger percentage of the product ammonia present.

The position of equilibrium for a reaction depends on several variables; these include concentration, temperature and pressure if gases are involved in the reaction. We can adjust the amounts of reactants and products at equilibrium by simply changing the reaction conditions. The trick is to know exactly what conditions to change and that is exactly what Henri Le Chatelier figured out.

In chemistry we often use the words system and surroundings when discussing a particular chemical reaction. The system is simply the reacting chemicals and the surroundings are the test tubes, beakers and indeed the whole universe- it's everything except the reacting chemicals! A closed system is one where no reactants or products can escape from the apparatus and nothing from the surroundings can get in either. Le Chatelier's Principle is the idea that a system (the reacting chemicals) at equilibrium will oppose any changes applied to it. We will use the Haber process to try and explain Le Chatelier's principle in more detail. The equation for the Haber process is shown again below:

Nitrogen_(g) + hydrogen_(g) ⇌ ammonia_(g)

N_2(g) + 3H_2(g) ⇌ 2NH₃_(g)

Now the forward reaction:

N_2(g) + 3H_2(g) → 2NH₃_(g)

is exothermic, that is it releases heat energy to the surroundings. The back reaction:

2NH_3(g) → N_2(g) + 3H_2(g)

is endothermic, that is it removes heat energy from the surroundings.

ammonia, nitrogen and hydrogen in flask to demonstrate how changes in temperature, pressure and concetration of gases will change the position of equilibrium.

So if you have a mixture of nitrogen, hydrogen and ammonia in a flask at equilibrium. The position of equilibrium for this reaction lies very much to the left and unfortunately there is very little ammonia in the flask. The challenge for chemists is how to adjust this equilibrium mixture so that the amount of the valuable product ammonia can be increased.

What would happen if we warmed the flask? Well according to Le Chatelier's principle the reaction should oppose any change you make to it; that is it will try and remove the heat you added. How would the added heat be removed? An endothermic reaction will remove heat! The only way to do this is to force the equilibrium position even more to the left, that is to increase the back reaction since this is an endothermic reaction:

2NH_3(g) → N_2(g) + 3H_2(g)

However this is bad news as it will reduce the amount of ammonia in the flask. So to increase the amount of ammonia instead of heating the flask, cool the mixture in the flask. Again according to Le Chatelier's Principle the system will try to generate heat; to oppose the cooling; so more nitrogen and hydrogen will react to make ammonia since the forward reaction is exothermic. However we know that cooling a reaction will slow it down, it will reduce the rates of the forward (R_f) and back reaction (R_b). So ultimately cooling it down will increase the amount of ammonia in the flask but it might take a very long time for the new equilibrium position in which there is more ammonia present to become established.
What would happen if you increased the pressure to the position of equilibrium in the Haber process? Well first thing is that pressure will only have an effect if there are gases involved in the reaction. So since all the chemical reacting here are gases then pressure will have an effect. Now you should know that at room temperature and pressure one mole of any gas will occupy 24 litres. So on the reactants side of the equation the total number of moles is 4 (3 moles of hydroegn and 1 mole of hydrogen) so this will occupy a volume of 96 litres. For the products we have 2 moles of gas; so this will occupy 48 litres at room temperature and normal pressure (standard temperature and pressure S.T.P). So obviously we can think of the reactants as the high pressure side of the equation and the products as the low pressure side. So if you increase the pressure on this equilibrium mixture of gases then it will force the equilibrium to the right, that is the low pressure side of the reaction. That is the equilibrium mixture will contain more ammonia. If you reduce the pressure then the equilibrium mixture will contain more reactants and less ammonia.
Similar arguments are true if you add more reactants or products then the system will no longer be at equilibrium and it will adjust the concentrations of all reactants and products to achieve a new equilibrium position. If you add more reactants it will force the equilibrium to the right, that is towards the products, similarly if you add more products the system will readjust to try and remove the added products by pushing the equilibrium towards the left hand side, that is the amount of reactants will increase.
How about adding a catalyst? Catalysts do not have any effect on the position of equilibrium; they will not affect the amount of reactants or products in a equilibrium mixture. What a catalyst will do is increase the rate of both the forward and reverse reactions in an equilibrium mixture. The catalyst will allow the system to achieve equilibrium faster.

Nitrogen dioxide and dinitrogen tetroxide

Dinitrogen tetroxide (N₂O₄) is colourless toxic gas with a very unpleasant smell. Despite the fact that it is a colourless gas it appears reddy brown since it exists in an equilibrium mixture with brown nitrogen dioxide gas. This can be shown as:

nitrogen dioxide_(g)⇌ dinitrogen tetroxide_(g)

2NO_2(g) ⇌ N₂O_4(g)

Flasks filled with nitrogen dioxide and dinitrogen tetroxide gas to show the effects of temperature on the position of equilibrium.

The forward reaction, the conversion of nitrogen dioxide into dinitrogen tetroxide is exothermic which means that the reverse or back reaction is endothermic.
What would happen to the position of equilibrium if we:

Placed a flask of this equilibrium mixture of gases in a beaker of iced water? Well according to Le Chatelier's principle the reaction should oppose any changes made to it. So in this case the change is a reduction in temperature so the position of equilibrium will shift to the right; that is increase the concentration of dinitrogen tetroxide since this reaction is exothermic and will generate heat energy. This would mean the colour of the equilibrium mixture of gases in the flask should fade, since the concentration of the brown gas nitrogen dioxide is reducing.

Using the opposite reasoning if we were to place a flask containing an equilibrium mixture of these two gases in a beaker of hot water, what would happen to the position of equilibrium? Well according to Le Chatelier's principle the system will try to oppose the temperature rise, so it will push the equilibrium in the direction of the endothermic reaction; that is there will be more nitrogen dioxide in the equilibrium mixture. This means the flask will get darker. This is shown in the image opposite.

Model equation showing the equilibrium reaction between nitrogen dioxide and nitrogen tetraoxide gas.

Le Chatelier's Principle and pressure

Using Le Chatelier's principle to explain changes in the position of equilibrium due to changes in pressure only applies to reactions involving gases. Solids and liquids are not affected by changes in pressure. As an example consider the reversible reaction that occurs between the colourless gas nitrogen tetroxide (N₂O₄) and the brown gas nitrogen dioxide (NO₂).

What would happen if an equilibrium mixture of these two gases was placed in a syringe and compressed? Well this would increase the pressure. So what would happen this time to the concentration of gases at equilibrium?

You can see from the equation opposite that 1 mole of the colourless N₂O₄ gas dissociates to form 2 moles of the brown NO₂ gas. Since there is 1 mole of gas on one side of the equation and 2 moles on the other we can say that the side with 1 mole of gas will be the low pressure side and the side with 2 moles of gas will be the high pressure side. The image below shows what happens to an equilibrium mixture of N₂O₄ and NO₂ gases in a syringe as the pressure is reduced and increased: Using syringes to show how the equilibrium mixture of nitrogen dioxide and nitrogen tetroxide is effected by changes in pressure

Using syringes to show how the equilibrium mixture of nitrogen dioxide and nitrogen tetroxide is effected by changes in pressure

Key Points

A reaction at equilibrium is one where the rates of the forward and reverse reactions are the same.
Le Chatelier's principle explains what happens to reactions at equilibrium. According to his principle a reaction at equilibrium will change or alter the position of equilibrium in order to minimise or cancel out the change applied.

Le Chatelier's principle